Covalent vs Ionic bonding
Covalent vs Ionic Bonding - Similarities and Application to Compounds Chemical bonding is typically taught in terms of four separate types: covalent, polar covalent, ionic and metallic. Whilst there are significant differences between the most extreme examples, bonding actually exists on a spectrum. Review of "Classical" Bonding Definitions Covalent - The equal sharing of electrons. Polar covalent - The unequal sharing of electrons. Ionic - An electrostatic interaction resulting from the transfer of electrons from one atom to another. Metallic - An electrostatic interaction resulting from the delocalisation of electrons. The Covalent-Ionic Spectrum Covalent and ionic bonding is often taught as being black or white, like these forms of bonding are two sides of the same coin. This is not an accurate way of representing chemical bonds. A more realistic approach is based on the electronegativity of the two atoms and involves considering a spectrum starting at a 100% covalent bond and ending at a 100% ionic bond. We will focus on bonds to chlorine for the sake of this discussion. ] A 100% covalent bond involves the equal sharing of electrons. This can only occur in homoatomic molecules, for example F2, O2, H2 and Cl2. The molecule is neutral, represented by the grey colour of the picture across. Even this is not a static picture as the electrons move around and can be influenced by neighbouring molecules ] As we increase the electronegativity difference between the two atoms, in this case by changing a chlorine atom for a hydrogen atom, we get uneven sharing of electrons. This leads to a partial positive charge on H (δ+) and a partial negative charge on Cl (δ-). These are represented by the red and blue colours respectively of the picture across. The bonding in HCl is roughly 82% covalent (and therefore roughly 18 % ionic, see below). We refer to this area of the spectrum as polar or polar covalent bonding. ] As we continue to increase the electronegativity difference further, in this case by changing a hydrogen atom for a sodium atom, the sharing becomes so uneven that it is more accurate to describe the bonding in terms of ions (i.e. ionic bonding). In other words sodium atom as loosened it's grip on the electron pair to such an extent that it appears to have transferred an electron to chlorine to form NaCl. This is represented in the diagram across. An important note is that, unlike the diagram, there is no such thing as a 100% ionic bond. There is always at least a bit of sharing between the atoms (so the red and blue areas should slightly overlap). The closest we have to a pure ionic species is CsF, which is roughly 91% ionic (and therefore ~9% covalent). In simplistic terms, we have seen that there are partial charges on the H and Cl atoms in HCl. As we know, opposite charges attract each other, therefore there must be an electrostatic component to the bonding (which we see as 18% ionic bonding). Bonding Triangles ] As we have established, there is a spectrum between 100% covalent and ~91% ionic with polar bonds in between. Metallic bonding can also be taken into account and we can see how these four bonding types relate to each other by using a bonding triangle (more accurately a van Arkel–Ketelaar triangle). These triangles take the extremes of ionic, covalent and metallic bonding and place them in the corners (see across). These plot the average electronegativity of an element against the electronegativity difference between two atoms. ] Here is an example of a bonding triangle with a variety of simple compounds. On this diagram covalent (and polar covalent) bonding is shown in yellow, ionic bonding in red and metallic bonding in blue. Inspection of this triangle shows how arbitrary lables such as "covalent" and "ionic" can be when dealing with the orange area at around the 50-50 ionic-covalent region. For example, according to this triangle, AlN sits exactly on the boundary so neither "ionic" nor "covalent" accurately describe the bonding in this compound. There is little difference between SiO2 and BeO yet one is called covalent and the other ionic. Classification of Compounds When attempting to classify simple compounds as "ionic" or "covalent" it is incredibly unhelpful to consider rules such as "compounds containing a metal and a non-metal are ionic" and "compounds containing only non-metals are covalent". That is simply not true. It may be a helpful guideline in exams at the start of your chemical education, but that does not make it correct! As previously mentioned, CsF is the most "ionic" compound that we can make and even that is not 100% ionic. 100% covalent molecules exist, but even they aren't 100% covalent 100% of the time It is far more accurate to consider the difference in electronegativities between the two atoms. There is a simple guideline as to whether a bond is covalent, polar or ionic based on the difference in electronegativities (ΔEN): pure covalent: ΔEN = 0 nonpolar covalent: 0 < ΔEN < 0.5 polar covalent: 0.5 < ΔEN < 1.7 ionic: ΔEN > 1.7 These are rough values, especially the polar/ionic interface. You'll see this number anywhere between 1.5 and 2 in different sources. Calculation of the Contribution of Ionic Character to a Chemical Bond ] There is a straight forward equation you can use to estimate the ionic contribution to a chemical bond in simple molecules. That equation is shown across. I is the percentage ionic character of a bond. μobs (pronounced "mu") is the dipole moment (D, Debyes) of the bond and μionic is the dipole moment of the bond if it was 100% ionic. μobs can be found in text books like Lange's handbook of chemistry (Table 5.18, page 5.130), whereas μionic you can calculate yourself. μionic is simply e x ''r where e is the charge of an electron and r is the bond length. If you use e in esu and r in Å (angstroms) the equation simply becomes μionic = 4.80320440 ''x r. Therefore I = (100 x'' μobs)/(4.803 ''x r). Bonding in f block complexes ] Lanthanide metals have low electronegativites (~1.1 - 1.3) but bonding in lanthanide complexes is particularly ionic in character. This is because, unlike most other elements, their valence orbitals (4f) are not really accessible, in other words they are "core-like". Covalent bonding requires orbital overlap otherwise electrons can not be shared. This is demonstrated in the image across. The 5s and 5p orbitals are lower in energy than the 4f orbitals and therefore get filled first. The graph (a radial distribution function) across in simplistic terms plots the probability of finding an electron against the distance from the nucleus. As 4f electrons are far more likely to be found closer to the nucleus than both the 5s and 5p (filled) orbitals. The 4f electrons are therefore not very accessible, causing lanthanide bonding to be predominately ionic in nature. 5f electrons are further from the nucleus so actinides sit between lanthanides and transition metal complexes in terms of bonding.